Bonding

1
Bonding

• Ch 8

2
Objectives

• SWBAT identify the bond type of a molecule by
  using electronegativity differences.
• SWBAT identify the polarity in small molecules.

3
Types of Bonding

• Ionic
• Covalent
• Metallic

4
Octet Rule

• Atoms tend to gain, lose or share electrons until
  they are surrounded by 8 valence electrons.

5

• Ionic bonding
• Transfer of electrons from the metal to the
  non-metal.
• Formation of cations and anions

Na Cl
Na Cl-
6
Ionic Bonding

• When an extremely electronegative atom, like
  fluorine, bonds with an electropositive atom,
  like sodium, the resulting bond is ionic due to
  the huge difference in electronegativity
  (difference gt 1.7).

7
Ionic Bonding

• The electronegative atom's pull on the bonding
  electrons is so strong that it pulls the bonding
  electron off the electropositive atom resulting
  in two oppositely charged ions which are held
  together by electrostatic attraction (an ionic
  bond).

8
Energetics of Ionic Bond Formation

• If you look in Appendix C (?Hf values) you will
  find that the heat of formation for ionic
  compounds is exothermic.
• The removal (or loss) of electrons is always an
  endothermic process.
• When a non-metal gains an electron the process is
  generally exothermic.

9
Lattice Energy

• The principal reason that ionic compounds are
  stable is the attraction between ions of unlike
  charge.
• Lattice energy is the energy required to
  completely separate a mole of a solid ionic
  compound into its gaseous ions.

10
Lattice Energy Problems

• Look at the example on BL page 267
• We will go through the example as a group.
• Try problems 19 and 20 at the end of the
  chapter.

11
Covalent Bonding

• Bonding between non-metals consists of two
  electrons shared between two atoms.
• In covalent bonding, the two electrons shared by
  the atoms are attracted to the nucleus of both
  atoms. Neither atom completely loses or gains
  electrons as in ionic bonding.
• http//www.elmhurst.edu/chm/vchembook/152Apolar.h
  tml

12
Types of Covalent Bonding

• Polar Covalent Bonding
• results when two different non-metals unequally
  share electrons between them.
• Non-Polar Covalent Bonding
• results when two identical non-metals equally
  share electrons between them.

13
Non-polar covalent bonds

• Identical non-metallic atoms have identical
  electronegativity
• For example in an H-H bond both H atoms have the
  same electronegativity so the bond is non-polar
• Different atoms can have the same
  electronegativity, such as, N-Cl.
• Both N and Cl have an electronegativity of 3.0 so
  the bonding electrons will be shared equally
  between the two atoms resulting in a non-polar
  covalent bond.

14
Polar Covalent Bonds

• When atoms of similar, but different,
  electronegativities (a difference lt 1.7) bond,
  the more electronegative atom has a greater share
  of the bonding electrons than the less
  electronegative atom.
• The more electronegative atom has a partial
  negative charge, and the less electronegative
  atom has a partial positive charge. The resulting
  covalent bond is called a polar covalent bond.

15
Show the Class

• Show how to draw a partial positive and partial
  negative charge symbol.

16
Multiple Bonds

• The length of the bond decreases with increasing
  numbers of bonds.

17
Comparison of Bonding Types
18
Compare Ionic, Polar and Non-Polar Bonding
19
Electronegativity

• Electronegativity is the relative tendency of a
  bonded atom to attract electrons to itself.
• An atom with extremely low electronegativity,
  like a Group I metal, is said to be
  electropositive since its tendency is to lose
  rather than to gain, or attract, electrons.
• Non-metals are more electronegative than metals.
• http//www.ausetute.com.au/bondpola.html

20
Electronegativity Values and Bond Types

• Electronegativity values are useful in
  determining if a bond is to be classified as
  non-polar covalent, polar covalent or ionic.
• What you should do is look only at the two atoms
  in a given bond. Calculate the difference between
  their electronegativity values. Only the absolute
  difference is important.
• http//dbhs.wvusd.k12.ca.us/webdocs/Bonding/Electr
  oneg-Bond-Polarity.html

21
Electronegativity and Associated Bond Type

• lt 0.5 non-polar covalent
• 0.5 1.9 polar covalent
• 2.0 or greater ionic
• Note To find the electronegativity values
  of many common elements, look at the chart
  printed on the bottom of your orbital
  diagram.
• there are some exceptions

22
Calculating Electronegativity Values

• Calculate the electronegativity and determine the
  bond type of an O-H bond
• O has an electronegativity of 3.5
• H has an electronegativity of 2.1
• The difference in electronegativity is
• 3.5 - 2.1 1.4
• 1.4 is less than 1.7, so the resulting bond is
  polar covalent.

23
AP Students Use This Slide

• Here are the rules
• If the electronegativity difference (usually
  called ?EN) is less than 0.5, then the bond is
  nonpolar covalent.
• If the ?EN is between 0.5 and 1.6, the bond is
  considered polar covalent
• If the ?EN is greater than 2.0, then the bond is
  ionic. That, of course, leaves us with a problem.
  What about the gap between 1.6 and 2.0? So,
  rule 4 is
• 4. If the ?EN is between 1.6 and 2.0 and if a
  metal is involved, then the bond is considered
  ionic. If only nonmetals are involved, the bond
  is considered polar covalent.
• http//dbhs.wvusd.k12.ca.us/webdocs/Bonding/Electr
  oneg-Bond-Polarity.html

24
AP Example

• Sodium bromide
• (formula NaBr ENNa 0.9, ENBr 2.8) has a
  DEN 1.9
• Hydrogen fluoride
• (formula HF ENH 2.1, ENF 4.0)
• has the same DEN.
• We use rule 4 to decide that NaBr has ionic
  bonds and that HF has a polar covalent bond in
  each HF molecule.

25
Electronegativity Trend

• Electronegativity decreases down a Group in the
  Periodic Table as the atomic radius and number of
  inner electron shells both increase.
• Electronegativity increases across a Period of
  the Periodic Table, in general, due to increasing
  nuclear charge and decreasing atomic radius.
• For the commonly encountered atoms in high school
  science, the order in decreasing
  electronegativity isF gt O gt N Cl gt Br gt C S
  I gt P H gt Si

26
Metallic Bonding

• Metallic bonding is characterized by a
• sea of electrons.

27
Dipole

• Polarity results from the uneven partial charge
  distribution between various atoms in a compound.
  
• Atoms, such as nitrogen, oxygen, and halogens,
  that are more electronegative have a tendency to
  have partial negative charges.
• Atoms, such as carbon and hydrogen, have a
  tendency to be more neutral or have partial
  positive charges.

28
Dipole

• Electrons in a polar covalent bond are unequally
  shared between the two bonded atoms, which
  results in partial positive and negative charges.
  The separation of the partial charges creates a
  dipole.
• The word dipole means two poles the separated
  partial positive and negative charges. A polar
  molecule results when a molecule contains polar
  bonds in an unsymmetrical arrangement.

29
Dipole

• Non-polar molecules are of two types.
• Molecules whose atoms have equal or nearly equal
  electronegativities have zero or very small
  dipole moments.
• A second type of nonpolar molecule has polar
  bonds, but the molecular geometry is symmetrical
  allowing the bond dipoles to cancel each other.

30
Dipole Moment

• The dipole moment is a measure of the unevenness,
  or lack of symmetry, of the charge distribution
  in a molecule.
• The mathematical definition of the dipole moment
  involves adding up the size of each charge in the
  molecule multiplied by the average distance that
  charge is from an arbitrary origin.
• www.chem.unsw.edu.au/.../dipolemoments.html

31
Dipole Moment

• symbol for dipole moment
• ?

32
Formal Charge

• See page 280 in the old B L text

http//www.chemprofessor.com/bonding_files/image03
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